Sulfur tetrafluoride

"SF4" redirects here. For the video game, see Street Fighter IV.
Sulfur tetrafluoride
IUPAC name
Sulfur(IV) fluoride
Other names
Sulfur tetrafluoride
7783-60-0 YesY
3D model (Jmol) Interactive image
ChEBI CHEBI:30495 YesY
ChemSpider 22961 YesY
ECHA InfoCard 100.029.103
PubChem 24555
RTECS number WT4800000
UN number 2418
Molar mass 108.07 g/mol
Appearance colorless gas
Density 1.95 g/cm3, 78 °C
Melting point 121.0 °C
Boiling point 38 °C
Vapor pressure 10.5 atm (22°C)[1]
Seesaw (C2v)
0.632 D[2]
Main hazards highly toxic
Safety data sheet ICSC 1456
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
US health exposure limits (NIOSH):
PEL (Permissible)
REL (Recommended)
C 0.1 ppm (0.4 mg/m3)[1]
IDLH (Immediate danger)
Related compounds
Other anions
Sulfur dichloride
Disulfur dibromide
Sulfur trifluoride
Other cations
Oxygen difluoride
Selenium tetrafluoride
Tellurium tetrafluoride
Polonium tetrafluoride
Related sulfur fluorides
Disulfur difluoride
Sulfur difluoride
Disulfur decafluoride
Sulfur hexafluoride
Related compounds
Thionyl fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Sulfur tetrafluoride is the chemical compound with the formula SF4. This species exists as a gas at standard conditions. It is a corrosive species that releases dangerous HF upon exposure to water or moisture. Despite these unwelcome characteristics, this compound is a useful reagent for the preparation of organofluorine compounds,[3] some of which are important in the pharmaceutical and specialty chemical industries.


Sulfur in SF4 is in the formal +4 oxidation state. Of sulfur's total of six valence electrons, two form a lone pair. The structure of SF4 can therefore be anticipated using the principles of VSEPR theory: it is a see-saw shape, with S at the center. One of the three equatorial positions is occupied by a nonbonding lone pair of electrons. Consequently, the molecule has two distinct types of F ligands, two axial and two equatorial. The relevant bond distances are S–Fax = 164.3 pm and S–Feq = 154.2 pm. It is typical for the axial ligands in hypervalent molecules to be bonded less strongly. In contrast to SF4, the related molecule SF6 has sulfur in the 6+ state, no valence electrons remain nonbonding on sulfur, hence the molecule adopts a highly symmetrical octahedral structure. Further contrasting with SF4, SF6 is extraordinarily inert chemically.

The 19F NMR spectrum of SF4 reveals only one signal, which indicates that the axial and equatorial F atom positions rapidly interconvert via pseudorotation.[4]

Intramolecular dynamic equilibration of SF4.

Synthesis and manufacture

SF4 is produced by the reaction of SCl2, Cl2, and NaF:

SCl2 + Cl2 + 4 NaF → SF4 + 4 NaCl

Treatment of SCl2 with NaF also affords SF4, not SF2. SF2 is unstable, it condenses with itself to form SF4 and SSF2.[5]

Use of SF4 for the synthesis of fluorocarbons

In organic synthesis, SF4 is used to convert COH and C=O groups into CF and CF2 groups, respectively.[6] Certain alcohols readily give the corresponding fluorocarbon. Ketones and aldehydes give geminal difluorides. The presence of protons alpha to the carbonyl leads to side reactions and diminished (30–40%) yield. Also diols can give cyclic sulfite esters, (RO)2SO. Carboxylic acids convert to trifluoromethyl derivatives. For example treatment of heptanoic acid with SF4 at 100-130 °C produces 1,1,1-trifluoroheptane. Hexafluoro-2-butyne can be similarly produced from acetylenedicarboxylic acid. The coproducts from these fluorinations, including unreacted SF4 together with SOF2 and SO2, are toxic but can be neutralized by their treatment with aqueous KOH.

The use of SF4 is being superseded in recent years by the more conveniently handled diethylaminosulfur trifluoride, Et2NSF3, "DAST", where Et = CH3CH2.[7] This reagent is prepared from SF4:[8]

SF4 + Me3SiNEt2 → Et2NSF3 + Me3SiF

Other reactions

Sulfur chloride pentafluoride (SF
), a useful source of the SF5 group, is prepared from SF4.[9]

Hydrolysis of SF4 gives sulfur dioxide:[10]

SF4 + 2 H2O → SO2 + 4 HF

This reaction proceeds via the intermediacy of thionyl fluoride, which usually does not interfere with the use of SF4 as a reagent.[5]


reacts inside the lungs with moisture:[11]

SF4 + 2 H2O → SO2 + 4 HF


  1. 1 2 3 4 "NIOSH Pocket Guide to Chemical Hazards #0580". National Institute for Occupational Safety and Health (NIOSH).
  2. Tolles, W. M.; W. M. Gwinn, W. D. (1962). "Structure and Dipole Moment for SF4". J. Chem. Phys. 36 (5): 1119–1121. doi:10.1063/1.1732702.
  3. C.-L. J. Wang, "Sulfur Tetrafluoride" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289.
  4. Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  5. 1 2 F. S. Fawcett, C. W. Tullock, "Sulfur (IV) Fluoride: (Sulfur Tetrafluoride)" Inorganic Syntheses, 1963, vol. 7, pp 119–124. doi:10.1002/9780470132388.ch33
  6. Hasek, W. R. "1,1,1-Trifluoroheptane". Org. Synth.; Coll. Vol., 5, p. 1082
  7. A. H. Fauq, "N,N-Diethylaminosulfur Trifluoride" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289.
  8. W. J. Middleton, E. M. Bingham. "Diethylaminosulfur Trifluoride". Org. Synth.; Coll. Vol., 6, p. 440
  9. Nyman, F., Roberts, H. L., Seaton, T. Inorganic Syntheses, 1966, Volume 8, p. 160 McGraw-Hill Book Company, Inc., 1966, doi:10.1002/9780470132395.ch42
  10. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
  11. Johnston, H. (2003). A Bridge not Attacked: Chemical Warfare Civilian Research During World War II. World Scientific. pp. 33–36. ISBN 981-238-153-8.
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